Chemical Bonding - Real-life applications

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Atoms, Electrons, and Ions

Today, chemical bonding is understood as the joining of atoms through
electromagnetic force. Before that understanding could be achieved,
however, scientists had to unlock the secret of the electromagnetic
interactions that take place within an atom.

The key to bonding is the electron, discovered in 1897 by English
physicist J. J. Thomson (1856-1940). Atomic structure in general, and
the properties of the electron in particular, are discussed at length
elsewhere in this volume. However, because these specifics are
critical to bonding, they will be presented here in the shortest
possible form.

At the center of an atom is a nucleus, consisting of protons, with a
positive electrical charge; and neutrons, which have no charge. These
form the bulk of the atom's mass, but they have little to do
with bonding. In fact, the neutron has nothing to do with it, while
the proton plays only a passive role, rather like a flower being
pollinated by a bee. The "bee" is the electron, and,
like a bee, it buzzes to and fro, carrying a powerful
"sting"—its negative electric charge, which
attracts it to the positively charged proton.

ELECTRONS AND IONS.

Though the electron weighs much, much less than a proton, it possesses
enough electric charge to counterbalance the positive charge of the
proton. All atoms have the same number of protons as electrons, and
hence the net electric charge is zero. However, as befits their highly
active role, electrons are capable of moving from one atom to another
under the proper circumstances. An atom that loses or acquires
electrons has an electric charge, and is called an ion.

The atom that has lost an electron or electrons becomes a positively
charged ion, or cation. On the other hand, an atom that gains an
electron or electrons becomes a negatively charged ion, or anion. As
we shall see, ionic bonds, such as those that join sodium and chlorine
atoms to form NaCl, or salt, are extremely powerful.

ELECTRON CONFIGURATION.

Even in covalent bonding, which does not involve ions, the
configurations of electrons in two atoms are highly important. The
basics of electron configuration are explained in the Electrons essay,
though even there, this information is presented with the statement
that the student should consult a chemistry textbook for a more
exhaustive explanation.

In the simplest possible terms, electron configuration refers to the
distribution of electrons at various positions in an atom. However,
because the behavior of electrons cannot be fully predicted, this
distribution can only be expressed in terms of probability. An
electron moving around the nucleus of an atom can be compared to a fly
buzzing around some form of attractant (e.g., food or a female fly, if
the moving fly is male) at the center of a sealed room. We can state
positively that the fly is in the room, and we can predict that he
will be most attracted to the center, but we can never predict his
location at any given moment.

As one moves along the periodic table of elements, electron
configurations become ever more complex. The reason is that with an
increase in atomic number, there is an increase in the energy levels
of atoms. This indicates a greater range of energies that electrons
can occupy, as well as a greater range of motion. Electrons occupying
the highest energy level in an atom are called valence electrons, and
these are the only ones involved in chemical bonding. By contrast, the
core electrons, or the ones closest to the nucleus, play no role in
the bonding of atoms.

Ionic and Covalent Bonds

THE GOAL OF EIGHT VALENCE ELECTRONS.

The above discussion of the atom, and the electron's place in
it, refers to much that was unknown at the time Thomson discovered the
electron. Protons were not discovered for several more years, and
neutrons several decades after that. Nonetheless, the electron proved
the key to solving the riddle of how substances bond, and not long
after Thomson's discovery, German chemist Richard Abegg
(1869-1910) suggested as much.

While studying noble gases, noted for their tendency not to bond,
Abegg discovered that these gases always have eight valence electrons.
His observation led to one of the most important principles of
chemical bonding: atoms bond in such a way that they achieve the
electron configuration of a noble gas. This has been shown to be the
case in most stable chemical compounds.

TWO DIFFERENT TYPES OF BONDS.

Perhaps, Abegg hypothesized, atoms combine with one another because
they exchange electrons in such a way that both end up with eight
valence electrons. This was an early model of ionic bonding, which
results from attractions between ions with opposite electric charges:
when they bond, these ions "complete" one another.

Ionic bonds, which occur when a metal bonds with a nonmetal are
extremely strong. As noted earlier, salt is an example of an ionic
bond: the metal sodium loses an electron, forming a cation; meanwhile,
the nonmetal chlorine gains the electron to become an anion. Their
ionic bond results from the attraction of opposite charges.

Ionic bonding, however, could not explain all types of chemical bonds
for the simple reason that not all compounds are ionic. A few years
after Abegg's death, American chemist Gilbert Newton Lewis
(1875-1946) discovered a very different type of bond, in which
nonionic compounds share electrons. The result, once again, is eight
valence electrons for each atom, but in this case, the nuclei of the
two atoms share electrons.

In ionic bonding, two ions start out with different charges and end up
forming a bond in which both have eight valence electrons. In the type
of bond Lewis described, a covalent bond, two atoms start out as atoms
do, with a net charge of zero. Each ends up possessing eight valence
electrons, but neither atom "owns" them; rather, they
share electrons.

LEWIS STRUCTURES.

In addition to discovering the concept of covalent bonding, Lewis
developed the Lewis structure, a means of showing schematically how
valence electrons are arranged among the atoms in a molecule. Also
known as the electron-dot system, Lewis structures represent the
valence electrons as dots surrounding the chemical symbols of the
atoms involved. These dots, which look rather like a colon, may be
above or below, or on either side of, the chemical symbol. (The dots
above or below the chemical symbol are side-by-side, like a colon
turned at a 90°-angle.)

To obtain the Lewis structure representing a chemical bond, it is
first necessary to know the number of valence electrons involved. One
pair of electrons is always placed between elements, indicating the
bond between them. Sometimes this pair of valence electrons is
symbolized by a dashed line, as in the system developed by Couper. The
remaining electrons are distributed according to the rules by which
specific elements bond.

MULTIPLE BONDS.

Hydrogen bonds according to what is known as the duet rule, meaning
that a hydrogen atom has only two valence electrons. In most other
elements—there are exceptions, but these will not be discussed
here—atoms end up with eight valence electrons, and thus are
said to follow the octet rule. If the bond is covalent, the total
number of valence electrons will not be a multiple of eight, however,
because the atoms share some electrons.

When carbon bonds to two oxygen atoms to form carbon dioxide (CO
2
), it is represented in the Couper system as O-C-O. The Lewis
structure also uses dashed lines, which stand for two valence
electrons shared between atoms. In this case, then, the dashed line to
the left of the carbon atom indicates a bond of two electrons with the
oxygen atom to the left, and the dashed line to the right of it
indicates a bond of two electrons with the oxygen atom on that side.

The non-bonding valence electrons in the oxygen atoms can be
represented by sets of two dots above, below, and on the outside of
each atom, for a total of six each. Combined with the two dots for the
electrons that bond them to carbon, this gives each oxygen atom a
total of eight valence electrons. So much for the oxygen atoms, but
something is wrong with the representation of the carbon atom, which,
up to this point, is shown only with four electrons surrounding it,
not eight.

In fact carbon in this particular configuration forms not a single
bond, but a double bond, which is represented by two dashed
lines—a symbol that looks like an equals sign. By showing the
double bonds joining the carbon atom to the two oxygen atoms on either
side, the carbon atom has the required number of eight valence
electrons. The carbon atom may also form a triple bond (represented by
three dashed lines, one above the other) with an oxygen atom, in which
case the oxygen atom would have only two other valence electrons.

Electronegativity and Polar Covalent Bonds

Today, chemists understand that most bonds are neither purely ionic
nor purely covalent; rather, there is a wide range of hybrids between
the two extremes. Credit for this discovery belongs to American
chemist Linus Pauling (1901-1994), who, in the 1930s, developed the
concept of electronegativity—the relative ability of an atom to
attract valence electrons.

Elements capable of bonding are assigned an electronegativity value
ranging from a minimum of 0.7 for cesium to a maximum of 4.0 for
fluorine. Fluorine is capable of bonding with some noble gases, which
do not bond with any other elements or each other. The greater the
electronegativity value, the greater the tendency of an element to
draw valence electrons to itself.

If fluorine and cesium bond, then, the bond would be purely ionic,
because the fluorine exerts so much more attraction for the valence
electrons. But if two elements have equal electronegativity
values—for instance, cobalt and silicon, both of which are
rated at 1.9—the bond is purely covalent. Most bonds, as stated
earlier, fall somewhere in between these two extremes.

POLAR COVALENT BONDING.

When substances of differing electronegativity values form a covalent
bond, this is described as polar covalent bonding. Sometimes these are
simply called "polar bonds," but that is not as
accurate: all ionic bonds, after all, are polar, due to the extreme
differences in electronegativity. The term "polar covalent
bond" is much more specific, describing a bond, for instance,
between hydrogen (2.1) and sulfur (2.6). Because sulfur has a slightly
greater electronegativity value, the valence electrons will be
slightly more attracted to the sulfur atom than to the hydrogen atom.

Another example of a polar covalent bond is the one that forms between
hydrogen and oxygen (3.5) to form H
2
O or water, which has a number of interesting properties. For
instance, the polar quality of a water molecule gives it a great
attraction for ions, and thus ionic substances such as salt dissolve
easily in water. "Pure" water from a mountain stream is
actually filled with traces of the rocks over which it has flowed. In
fact, water—sometimes called the "universal
solvent"—is almost impossible to find in pure form,
except when it is purified in a laboratory.

By contrast, molecules of petroleum (CH
2
) tend to be nonpolar, because carbon and hydrogen have almost
identical electronegativity values—2.5 and 2.1 respectively.
Thus, an oil molecule offers no electric charge to bond it with a
water molecule, and for this reason, oil and water do not mix. It is a
good thing that water molecules attract each other so strongly,
because this means that a great amount of energy is required to change
water from a liquid to a gas. If this were not so, the oceans and
rivers would vaporize, and life on Earth could not exist as it does.

BOND ENERGY.

The last two paragraphs allude to attractions between molecules, which
is not the same as (nor is it as strong as) the attraction between
atoms within a molecule. In fact, the bond energy—the energy
required to pull apart the atoms in a chemical bond—is low for
water. This is due to the presence of hydrogen atoms, with their two
(rather than eight) valence electrons. It is thus relatively easy to
separate water into its constituent parts of hydrogen and oxygen,
through a process known as electrolysis.

Covalent bonds that involve hydrogen are among the weakest bonds
between atoms. (Again, this is different from bonds between
molecules.) Stronger than hydrogen bonds are regular, octet-rule
covalent bonds: as one might expect, double covalent bonds are
stronger than single ones, and triple covalent bonds are stronger
still. Strongest of all are ionic bonds, involved in the bonding of a
metal to a metal, or a metal to a nonmetal, as in salt. The strength
of the bond energy in salt is reflected by its boiling point of
1,472°F (800°C), much higher than that of water, at
212°F (100°C).